F. *J. DILLEMUTH, S.J., D. R. SKIDMORE, S.J., AND C. C. SCHUBERT, S..J.
1496
complexes, it appears that the reduction of Pu(V1) by Sn(I1) may offer an opportunity to obtain the entropy of an activated complex with a zero or negative charge. Acknowledgments.-The author wishes to express his appreciation to Mr. R. H. Moore for the
Vol. 64
preparation of the IBM-704 program used in the evaluation of the specific rate constants mid for helpful discussions on problems of statistical interest. Also, helpful discussions with Drs. J. F. Lemons, C. E. Holley, Jr., and T. W. Newton concerning this research are gratefully acknowledged.
THE REACTIOY OF OZOKE WITH METHANE BY FREDERICK J. DILLEMUTH, S.J., DUANER. SKIDMORE, S.J. ASD CLARENCE C. SCHUBERT, S.J. Department of Chrnistry, Fordham Univmity, New York 68, New York Receaved April 1 4 , 1060
The reaction of ozone with methane was studied in temperature controlled, infrared, gas cells. Infrared analysis showed the major reaction products to be carbon monoxide, carbon dioxide, formic acid and water vapor. It was found that the reaction was substantially independent of Pyrev and sodium chloride Rurfaces. The activation energy calculated on the basis of the equation -d[03]/dt = k[Oa][CH4]was found to be 15.3 kcal./mole m t h oxygen as the ozone diluent. The preexponential term was 1.6 X 10” cc. mo1e-I set.-*. When oxygen was excluded from the reaction mixture of methane and ozone, a value for the activation energy of 13.9 kcal. was found and a pre-exponential term 1.4 X 1Olo cc. mole-’ sec.-l, a value regarded as being not significantly different. The low activation energy found for this reaction is taken as a strong indication that, the reaction mechanism does not include a prior decomposition of ozone followed by an attack of oxygen atoms on methane. It was found that there is an approximate equivalence of gram atoms of oxygen fixed in products and moles of ozone consumed.
I. Introduction I n 1898, Otto’ first studied the reaction of methane arid ozone both at room temperature and at looo,obsening the formation of aldehyde. Blair and Wheeler,2 repeating the same experiment at loo”, noted that 53% of the ozone present had reacted with methane after two minutes. No carbon monoxide was detectable. The effect of surface on the reaction was little, if any, serving only t o destroy some of the ozone. Briner and C a r ~ e l l e rreporting ,~ the action of ozone in the slow combustion of propane and butane, postulate a chain mechanism in which the chain initiating step is hydrogen abstraction by a molecule of ozone RCE3
+ 0s +RCH + HzO +
0 2
(I)
Schubert and Pease,4 in their studies on the oxidation of the lower paraffin hydrocarbons, found that the reaction of methane with ozonized oxygen occurred with an activation energy of 14,900 cal./ mole. To explain this relatively low activation energy, it was postulated that the reaction might occur between the hydrocarbon and an ozone molecule in a triplet low-lying excited electronic state. Mora recently, Kleimenov, et U Z . , ~ using a flow system reinvestigated the methane-ozonized oxygen reaction. Because oxygenated products appeared only at and above the temperature at which ozone was found to decompose, these investigators concluded that the oxidation of methane and other hydrcicarbons is primarily effected by an atomic: mechanism, viz. 0 3
--R H
+0
2
+0
+ 0 +R + .OH etc.
(11) (111)
(I) M M. Otto, Ann. chtm. phys , [VII] is, 109 (1898). . Ind. (London),41, (2) E. W. Blair and T. S. Wheeler, J . S O C Chem. 303 (1922). (3) E. Briner and J. Carceller, Helr. Chzm. Acto, 18, 973 (1935). (4) C. C. Schiibert. 9. J. and R. N. Pease, J . Am. Chem. Soc., 7 8 , 2044 (1956). (5) N.A. Kleimenov, I. N. Antonova, A. M. Markevich and -4. B. Nabaldiam, Zhur. Far.Khzm.. X X X , 794 (1956).
Reported here are the results of studies concerning the reaction of ozone with methane, both in the presence and in the absence of added oxygen. 11. Experimental Tank oxygen (Ohio Chemical and Surgical Company,
U.S.P.) was dried over calcium chloride and silica gel. It was next ozonized in a Siemens-type ozonizer, the effluent ozone-oxygen mixture being ca. 3 mole %in ozone. Ozone-
oxygen mixtures of concentrations higher than 3 mole % in ozone, were prepared according to a method outlined by Cook.6 The ozone-oxygen mixture was_ allowed to pass through a silica gel column cooled to - ( S o with a slurry of acetone and carbon dioxide snow. At this temperature, ozone adsorbed on the silica gel, while oxygen passed through the column unabsorbed to any significant extent. The effluent oxygen, after a washing with 2% potassium iodide solution in order to destroy any unadsorhed ozone, was allowed to escape into the atmosphere. The concentrated ozone was adsorbed as a wide purple band on the silica gel column. When required for use, conccntrated ozone was desorbed from the silica gel a t room temperature in a stream of dried oxygen or nitrogen. Methane (hlatheson Co., C.P.) was purified from traces of carbon dioxide by passing the gas through moist potassium hydroxide pellets and dried over calcium chloride. Methane and the enriched ozone were mixed in a capillary mixer and allowed to fill a previously evacuated ten cm. infrared gas absorption cell until the total pressure was one atmosphere. All gases were metered by calibrated flow meters. The oxidations were followed by an infrared spectrophotometric method of a n a l y ~ i s using ,~ a Perkin-Elmer Model 21 double beam spectrophotometer. The infrared absorbance of the various reactants and products was noted a t specific intervals during the course of the reaction. The ozone absorbance was measured a t 1055 cm.-l, carbon monoxide a t 2183 cm.-l, carbon dioxide a t 2344 cm.-l and formic acid at 1740 crn.-l. Concentrations of reactants and products were calculated by comparison with calibration curves constructed from reference spectra run on pure compounds a t various known partial pressures, under conditions as similar as possible to those obtaining during the oxidations. Methane concentrations could not be determined spectrophotometrically. Initial methane concentrations were obtained from initial flov rates. (6) G. A. Cook, A. D. Kiffer, C. V. Klumpp, .I. 13. hfalik and L. A. Spence, “Separation of Onone from Oxygen by a Sorption Process,” Tonawanda Research Laboratory, Linde Air Products Company, Tonawanda, New York.
Oct., 1960
THEREACTION OF OZONE WITH METHANE
The ten-cm. infrared cells, constructed of Pyrex glass, were fitted with sodium chloride windows. They were equipped with a water jacket for temperature control to within =k 0.25'. In double beam operation, the reference cell was filled up to one atmosphere pressure with either dried oxygen or nitrogen, depending on whichever gas had been used t o desorb the ozone from the silica gel. All stopcocks were lubricated with an ozone-resistant Halocarbon grease. Time zero was recorded from the moment when the reaction cell was brought up to temperature. Preliminary experiments showed that any reaction at room temperature which occurred in the brief time elapsing between filling the cells and the moment when the desired temperature was reached, was negligible.
111. Experimental Results (a) Products of the Ozone-Methane Reaction.-The infrared absorption spectra of the products of the reaction showed principally, carbon monoxide, carbon dioxide, formic acid and water vapor. Although a slight amount of unreacted ozone could still be observed when the reaction was essentially complete, its concentration was usually of the order of 0.1 mmole/l., and product ahsorptions interfered to such an extent that any quantitative estimate of ozone concentration beyond this point was unreliable. Infrared bands characteristic of higher paraffin hydrocarbons, arising, presumably, from the reaction R. .R' + RR', were not observed; neither mas the presence of such compounds considered likely.' Formaldehyde was not detected in any of the infrared spectra, though trace amounts were probably present. Figure 1 shows a representative run conducted with ozonized oxygen at 66.6". In the whole series of runs, the hydrocarbon concentration was, on the average, 2.91 X 10-2 mole/l., while the ozone mixture was held a t apmole/l. In this manner proximately 1.55 X a gas mixture rich in hydrocarbon was always assured. This experimental condition was forcibly suggested by a severe explosion which shattered one of the gas absorption cells in a preliminary investigation. It was observed that in all cases the concentration of carbon monoxide produced was roughly 1.5-2 times as great as that of the carbon dioxide. Formic acid appeared in all of the runs, but its concentration in no case exceeded 1.2 X mole/l. Conceivably, in the runs a t lower temperatures, some of the acid condensed on the cell walls and thus escaped detection. Product yields, determined by infrared analysis for several runs, both when oxygen was added and when it was excluded from the reaction mixture, are given in Table I. It is not felt that the results are too significantly different in both cases. The calculated values for the amounts of water produced are based on the assumption that two moles of water are produced for every mole of carbon appearing in products. Since the infrared spectra showed no hydrogen-containing oxidation products other than water, except in trace amounts, this would appear to be a reasonable approach toward a product balance. The ratio of gram atoms of oxygen fixed in product to moles of ozone decomposed was calculated on the basis of sixteen independent runs. I n
1.0
c
s X
2 .$ 0.5 E
+
(7) E. W. Steacie and N. A. Parlee. Con. J . Research. B16, 203 (1 988).
1497
I
0
0
--
2
FORMIC ACID
I
I
I
I
I
4
6
8
10
12
Seconds x IO-'. Fig. 1.-Methane plus ozone a t 66.6'.
TABLE I MOLES O F PRODUCT P E R
Run no.
T:mp.,
C.
COI
MOLEO F OZONE CONSUMED
co
HtO
209 162 225 164
35 43.9 45.3 64.4
Oxygen diluent 0.16 0.24 .360 .I30 .09 .22 .019 .308
167
35.4 44.4 48.4 64.4
Nitrogen diluent 0.080 0.184 0.53 .I22 .227 .70 .IO7 .223 -66 ,062 ,231 .587
163 185 166
0.81 .49 .63 ,658
Gram atoms oxygen fixed
1.37 1.11 1.03 1.10
0.874 1.17 1.09 0.943
eight of these oxygen was used as the ozone diluent. In the others oxygen was excluded by using nitrogen as the cariier gas for the ozone. The average showed that in the former case 0.997 gram atoms of oxygen mere fixed in product per mole of ozone consumed. I n the latter case, 1.02 gram atoms n-ere fixed per mole of ozone. This result has been taken to indicate that one mole of ozone reacts with one mole of the hydrocarbon with the ultimate formation of oxygenated products plus one mole of molecular oxygen. (b) Effect of Surface.-Since the reaction cell was composed of Pyrex glass and fitted with rock salt windows, it was felt desirable to investigate the effects of added Pyrex and salt surfaces on the course of the reaction independently. The former was accomplished by placing five lengths of 8 mm. diameter Pyrex tubing lengthwise into the cell, changing the Pyrex surface to volume ratio from 2.02 t o 5.79 ern.". Although the optical balance
1498
E". J . DILLEMUTH, S.J.,
n. R. SKIDMORE, S.J., AND c. c. SCHVRERT, S.J.
350 c. 43.90 x 57.3O 66.6"
0
b
'\X
1.5
' 0 .
Vol. 64
1 and 2. Since, in each series of experiments, the concentration of methane was far in excess of that of ozone, the hydrocarbon concentration was taken as effectively constant during the course of each run. The activation energy for the reaction was determined from Arrhenius' plots, both for the case when oxygen was used as the ozone diluent and when nitrogen was so employed. These values are given in Table 11, together with the rate constants calculated for 25" and the pre-exponential term A determined from the equation
0"
k = A exp( -E*/RT) cc. mole-' sm-1
$1.0 A
TABLE I1
0.5 -
X
*\
i \
-
X
kra (calcd.) (cc. mole-1 sec. -1)
E* (cal./moleI
A (CC.mole-' aec.-1)
15,350
Oxygen added 1.63 X 10"
0.82
13,900
Oxygen excluded 1.43 X IO'"
0.84
The variation in the activation energy and the pre-exponential term when oxygen was excluded from the reaction mixture, is not regarded as significant. 0 2 4 6 8 1 0 1 2 1 4 Discussion Seconds x 10-3. In all the experiments there was, on the average, Fig. 2.-Methane plus ozonized oxygen. an approximate equivalence between the number of the spectrophotometer ha;d to be modified of gram atoms of oxygen fixed in products and the somewhat to compensate for the interference of- moles of ozone disappearing. On the basis of this fered by the tubing to the passage of the infrared finding, a chain mechanism involving oxygen radiatio.n, the greater percentage of the radiation molecules has been excluded. Were a molecular was unhindered in its travel through the cell. The oxygen chain reaction important in the low teminsertion of a number of regular sodium chloride perature reaction of ozone with methane, there crystals into the body of the cell changed the salt would be expected a significantly greater amount of oxygenated products. surface t o volume ratio from 0.209 to 0.891 The results of this work show that the reaction It was found that in neit.her case were the kinetics of the reaction changed, nor did the use of nitrogen of ozone and methane, both with and without added rather than oxygen as the ozone diluent cause any oxygen, proceeds with an activation energy of apsignificant differences in this series of experiments. prox. 14-15 kcal. 'mole. The thermal decomposition of ozone, on the other hand, occurs, according (c) Kinetics.-In Fig. 2 the logarit'hm of ozone with an activation to recent investigations, lo partial pressure is plotted against time. From the linear piot's obtained for the reaction of ozone with eiiergy of 24.2 f 0.2 kcal./mole. It might be methane a t four temperatures, it was inferred t,hat questioned, therefore, in the light of this divergence the reaction :is first order with respect to ozone. of artirntion energy values, whether a mechanism Points corresponding to an ozone partial pressure of which invokes oxygen atoms and not the ozone less than 4 mm. were unreliable because product molecule itself as the initial attacking species, absorptions interfered with the fundamental ozone adequately fits the observed data. Consequently, band a t 1055 cm.-'. In general, marked deviat'ions it is not immediately obvious how the conclusions from linearity were observed at low ozone concen- drawn by Kleimenov,5 who postulates the prior decomposition of the ozone molecule and a subtrations.. sequent reaction betn-een methane and an oxygen From prev.ious on t,he ozonolysis of hydrocarbons, it was felt that the reaction could be atom, can be reconciled with the results of this work. With respect to the reaction rate constants, actitaken at3 second order, and the rate expression be vation energy and the pre-exponential term for the given as reaction of methane with ozone, we are impressed -d[Ojl/dt = k[Osl [CHd by the fact that, for all practical purposes, they where k, in this case, has the units cc. mole-' sec.-l. arc the same both in the presence and in the abData were talien from plots such as those of Figs. sence of added oxygen, arid. at least in the tempera(8) R. D. Cadle, "Chemical Reactions in Los Angeles Smog," Proceedings of the Second National Air Pollution symposium, Stanford, Calif., Stanford Research Institute, 1952, p. 31. (9) R. H . Sawye: and F. W. Behnke, Bbstracts of thr 132th Meeting of the American Chemical Society, Boston, Mass., April, 1959, p. 28-B.
(10)
S.'&- Benson and -4 E. Axivnrthy, Jr., J . Chem. Phgls., %6, 1718
(1957). (11) J. A. Zaslowsky. H. B. Crbach. F. Leighton and R. J. Wnuk, 4bstraots of t h e 135th 3Ieetina of t l i c .Inierlcan Cheroical Society, Boston, Mass., p . 23-R.
O p t . , 1960
OSMOTIC -4ND -ACTIVITY COEFFICIENTS OF SULFONIC ACIDS
ture range employed, independent of surface. The reaction is believed, therefore, to proceed by a similar mechanism in all the above cases. Acknowledgment.-This research was supported by a grant from the Petroleum Research Fund, administered by the American Chemical Society.
1499
Grateful acknowledgment is hereby made t o the donors of said fund. The authors are also indebted to Dr. Raymond R. Sawyer, Mr. Fred Behnke and the Perkin-Elmer Corporation of n'orwalk, Connecticut for the generous loan of long path infrared cells used in the course of this work.
THE EFFECT OF STRUCTURE O S THE OSMOTIC ASD ACTIVITI? COEFFICIENTS OF SOME SULFOXIC ACIDS AKD THEIR SALTS BY 0. D. BONNER'AND 0. C. ROGERS Department of Chemistry, University of South Carolina, Columbia, South Carolina Receaued Apral 16, 1960
Osmotic and activity coefficients are reported for benzenesulfonic and mesitylenesulfonic acids and the lithium and sodium salts of benzenesulfonic, p-ethylbenzenesulfonic, 2,5-dimethylbenzenesnlfonicand mesitylenesulfonic acids a t 25'. These measurements were made by isopiestic comparison of solutions of the sulfonates with sodium chloride solutions as reference standards. The osmotic and activity coefficients of the acids and salts are observed to decrease with increasing molecular weight. The magnitude of the values of these coefficients have the order Li > Xa > H for the higher molecular weight aulfonates, thus indicating that these sulfonic acids are probably only moderatelv strong acids.
Since the introduction of sulfonic acid type of cation-exchange resins which consist of sulfonate exchange sites on polystyrene-divinylbenzene matrices, there have been many attempts to interpret the equilibrium data obtained for exchange reactions involving various pairs of ions. At least two2 of these attempts have iiivolved the use of model compounds similar in structure to the ion exchanges. This work has been handicapped by the scarcity of data for the various aromatic sulfonic acids and their salts. Robinson and Stokes3 have reported osmotic and activity coefficients for the lithium, sodium and potassium salts of p tolucnesulfonic acid while Bonner2b and coworkers have reported data for p-toluenesulfonic, p-ethylbmzenesulfonic, 2,5-dimethylbenzenesulfoiiic, 4,I'-bibenzyldisulfonic and m-benzenedisulfoiiic acids. The only compound for which osmotic and activity coefficient data are available for both the parent acid and some of its salts is therefore p-toluenesulforiic acid. Experimental The various sulfonic acids with the exception of benzenesulfonic acid were prppared by sulfonation of the corresponding purified hydrocarbons. Benzenesulfonic acid was prepared by the hrdrolysis of beiizenesiilfonyl chloride. The sulfonates were rerowred f i orn the sulfuric arid solution by neutralization with potassiiim hydroxide in the case of bmrid :tiid sodium hpdro\ide for the other solulis were then recrystallized at lcwt three times from water or wntrr-methanol solutions and dried to consLant weight. S o structural isomers are possible for any of the sulfonates except the ethylbenzenesulfonates. The purity of the various salts was checked by passing solutions of weighed quantities of each salt through an ion-exchange column in the acid (hydrogen) form and titrating the liberated acid (1) The d a m repoited in this paper are from a project supported by the Vnited States Atomic Energy Commission. (2) (a) G . E. Myera and G. E. Boyd, THISJOURNAL, 60, 521 (1956); (b) 0. D. Bonner, V. E'. Holland and Linda Lou Smith, ibid., 6 0 , 1102
